1. 1.5 Oxidation and Reduction

2. Learning Outcomes Introduction to oxidation and reduction: simple examples only, e.g. Na with Cl 2, Mg with O 2, Zn with Cu 2+. Oxidation and reduction in terms of loss and gain of electrons. Oxidising and reducing agents. The electrochemical series as a series of metals arranged in order of their ability to be oxidised (reactions, other than displacement reactions, not required). Electrolysis of (i) copper sulfate solution with copper electrodes and (ii) acidified water with inert electrodes. (Half equations only required.)

3. Oxidation and reduction Oxidation = addition of oxygen to a substance C + O 2  CO 2 Reduction is loss of oxygen or addition of hydrogen CuO + H 2  Cu + H 2 O

4. examples Sodium + chlorine  sodium chloride Na + Cl  Na + + Cl - Na loses an electron [oxidised] Cl gains an electron [reduced]

6. Example 2 Magnesium +oxygen  magnesium oxide Mg + O  MgO Mg  Mg +2 loses 2 electrons [oxidation] O  O -2 gains 2 electrons [reduction]

7. Example 3 Zinc +copper sulphate  Zinc sulphate+ Copper Zn + Cu +2  Zn +2 + Cu Zinc loses electrons (oxidised) Copper gains electrons (reduced)

8. Oxidising agent A substance that causes oxidation in another substance

9. Reducing agent A substance that causes reduction in another substance.

10. Oxidation is loss of electrons; Reduction is gain of electrons CuO + H 2  Cu + H 2 O CuO  Cu +2 and O -2 Cu +2  Cu [gains 2 electrons] reduced H 2  H 2 +2 [loses 2 electrons] oxidised O -2  O -2 [ no change]

11. Oxidation numbers The charge that an atom has or appears to have assuming that the compound is ionic. Electrons always go the the most electronegative element

12. Oxidation number rules 1 Elements on their own = 0 H 2 = 0 Zn = 0 Cl 2 = 0

13. Oxidation number rules 2 Ions = same as charge Cu +2 = +2 O -2 = -2 Cl -1 = -1

14. Oxidation number rules 3 Charges of all elements in a compound = 0 CuSO 4 Cu = +2 S = +6 O 4 = -8 [O = -2] Total = +2 +6 –8 = 0

15. Oxidation number rules 4 Oxygen = -2 Exceptions are peroxides O = -1 [H 2 O 2, Na 2 O 2 ] OF 2 O = +2, F = -1

16. Oxidation number rules 5 Hydrogen = +1 Exceptions are the metal hydrides NaH Na = +1, H = -1

17. Oxidation number rules 6 Halogens [ Cl, F, I, Br] are always –1 except when joined to more electropositice element Cl 2 O Cl = +1, O = -2

18. Oxidation number rules 7 In a complex ion the sum of all the charges = the chartge on the ion. SO 4 - 2 S = +6, O 4 = -8 [O = -2] +6 –8 = -2

19. redox Oxidation is an increase on oxidation number Reduction is a decrease in oxidation number.

20. Electrochemical Series Electrochemical Series – Elements listed in order of ability to be oxidised

21. Metals King [K] Neptune [Na] Caught [Ca] Many [Mg] Angry [Al] Zulus [Zn] Fighting [Fe] Police [Pb] Constables [Cu] Having [Hg] Asthma [Ag] Attacks [Au]

22. Metals above hydrogen in the Reactivity Series react with acids to produce hydrogen gas. Zinc Potassium Sodium

23. Displacement of metals Displacement reactions occur when a metal from the electrochemical series is mixed with the ions of a metal lower down in the series. The atoms of the more reactive metal push their electrons on to ions of the less reactive metal.

24. Displacement More reactive metal displaces less reactive from a solution Mg + CuSO 4 = MgSO 4 + Cu Mg + Cu +2  Mg +2 + Cu Mg loses electrons (Oxidised) Cu +2 gains electrons (reduced)

25. Learning Outcomes Rusting of iron. Swimming-pool water treatment. Use of scrap iron to extract copper. Electroplating. Purification of copper. Chrome and nickel plating. Cutlery.

26. Rust Rust is the formation of iron oxides (usually red oxides), formed by the reaction of iron and oxygen in the presence of water or air moisture. Oxidation

27. Swimming pools The water in swimming pools is kept sterile by the addition of oxidizing agents, chlorine or chlorine compounds, which kill microorganisms by oxidation. The active agent is usually chloric(1) acid (HOCl). It may be formed in two ways 1. Direct chlorination of the water: Cl 2(aq) +H 2 O (l)  HOCl (aq) + Cl − (aq) + H + (aq) Note that when the Cl 2 reacts with the water it is both oxidized and reduced

28. Swimming pools 2. The addition of sodium chlorate(I) [sodium hypochlorite]: NaOCl (s) + H 2 O (l) Na + (aq) + OH − (aq) + HOCl (aq) Nowadays chlorine is not used, mainly on grounds of safety. Pools are sterilized with chlorine compounds, which produce chloric(I) acid when they dissolve in water. These compounds act in essentially the same way as chlorine. Sodium chlorate(I) is one such compound.

29. Use of scrap iron to extract copper. (Dissolved CuSO 4 ) + (Metallic Fe) ==> (Dissolved FeSO 4 ) + (Metallic Cu)

30. Electrolysis Chemical reaction caused by the passage of an electric current through a liquid known as the electrolyte

31. Definitions Electrolyte - liquid in which electrolysis takes place. Usually an ionic solution but it can also be a fused [melted] ionic compound Anode - positive electrode. Positive because the battery sucks electrons out of it Cathode. Negative electrode. Negative because the battery pumps electrons into it. Anion - negative ion. Called anion because it is attracted to the opposite charge of the anode Cation - positive ion. Called cation because it is attracted to the opposite charge of the cathode. Inert Electrodes - do not react with the electrolyte Graphite and Pt Active electrodes - react with electrolyte e.g. Copper and iron

32. Electrolysis

33. Electroplating Covering cathode in metal e.g. Cu by making it cathode in copper sulphate solution

34. Copper plating

35. Copper Plating Anode reaction Cu (s) = Cu 2+ (aq) + 2e - Anode loses mass as copper dissolves off Impurities [Au, Ag, Pt etc.] fall to bottom Cathode reaction Cu 2+ (aq) + 2e - = Cu (s) Cathode gains mass as Cu is deposited on it Cu is 99.9% pure

36. Learning Outcomes Mandatory experiment 1.2 (half equations only required, e.g. 2Br – – 2e – → Br 2 ). Demonstration of ionic movement. Demonstration of electrolysis of aqueous sodium sulfate (using universal indicator) and of aqueous potassium iodide (using phenolphthalein indicator) with inert electrodes. (Half equations only required.)

37. Ionic Movement During electrolysis of a solution of Copper Chromate in dil. Hydrochloric acid, positive ions (cations) are attracted to the negative electrode (cathode) and negative ions (anions) are attracted to the positive electrode (anode). If these ions are coloured, their movement may be observed visually. Examples of coloured ions include; copper(II) [Cu 2+ ] - blue chromate(VI) [CrO 42- ] – yellow

38. Q & A to Ionic Movement Expt (1) What colour is the copper(II) chromate solution? Copper(II) chromate solution is an olive green colour. (2) What colour is observed at the positive electrode after the power supply has been turned on for some time? A yellow colour is observed at the positive electrode. (3) What colour is observed at the negative electrode after the power supply has been turned on for some time? A blue colour is observed at the negative electrode. (4) Explain in terms of the movement of ions why different colours are formed at each electrode. When the circuit is completed, positive copper ions (Cu 2+ ) are attracted to the negative electrode. These ions have a blue colour. Similarly negative chromate(VI) ions (CrO 42- ) are attracted to the positive electrode. These ions are coloured orange. (5) What is the function of the dilute hydrochloric acid? The dilute hydrochloric acid is required to complete the circuit.

39. Electrolysis of Sodium sulphate Solution of Na 2 SO 4 + universal indicator H + ions are produced at the positive electrode (oxidation of O 2- in water) while OH - ions are produced at the negative electrode as the H + in water is reduced to H 2 (g).

40. Sodium Sulphate and Universal Indicator

41. Electrolysis of Sodium Sulphate Red is acid at the positive electrode 2H 2 O(l)  O 2 (g) + 4H + (aq) + 4 e - lose electrons = oxidation = anode Purple is base at the negative electrode H 2 O(l) + 2 e -  H 2 (g) + 2OH - (aq) gain electrons = reduction = cathode

42. Electrolysis of Potassium Iodide Solution of KI + phenolphthalein Brown I 2 (s) forms at the positive electrode and some yellow/orange I 3 - forms in solution. At the negative electrode, H + is again reduced to H 2 (g) and the phenolphthalein turns pink due to the OH - ions.

43. Electrolysis of Potassium Iodide

44. KI  K + + I- Iodide loses electrons  Brown iodine 2I-  I 2 + 2e - Anode, Oxidation H 2 O  H + + OH - OH - is basic, Phenolphthalein Purple.